The Periodic Table is trendy! Episode 14 recaps Coulomb’s Law (1:10) and describes how it applies to atomic structure (2:00).
The Periodic Table is trendy! Episode 14 recaps Coulomb’s Law (1:10) and describes how it applies to atomic structure (2:00). We can use Coulomb’s Law to predict atomic radii (3:30) ionic radii (4:00). The episode also defines ionization energy, electronegativity and electron affinity (4:45), describes their periodic trends and ties them to Coulomb’s Law and atomic structure (5:33). But a trend is not always followed. Therefore we also recap two exceptions: the ionization energy of oxygen and nitrogen (6:23) as well as the first ionization energy of boron and beryllium (6:50). And because it is so important, we also have a test taking tip for you (7:28)!
Question: Removing a second electron and forming a 2+ cation, requires ____ energy than removing the first electron. (8:31)
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Hi and welcome to the APsolute Recap: Chemistry Edition. Today’s episode will recap Coulomb’s Law and Periodic Trends
Lets Zoom Out:
Unit 1 - Atomic Structure and Properties
Topic - 1.5 and 1.7
Big idea - Structure and Properties
Have you heard about the latest trend? What trend? There is always a trend: what people wear, what apps they use, what games are being played. What does it mean when we say something is “trending”? Usually that indicates that an increasing number of people do it. Not everyone, but a lot of people. We even use a word with a positive connotation: Being trendy. Uuuuhhh! But not only you and your friends are trendy - or might decide on purpose NOT to follow a trend. The periodic table is also trendy! Today’s episode recaps periodic trends and Coulomb’s Law.
Let’s zoom in:
Coulomb’s Law states that opposite charges attract and like charges repel. The strength of the force is determined by the charge as well as the distance. Let’s use a magnet as an example: If you compare two small kitchen magnets with large industrial magnets, the forces of attraction and repulsion are different: smaller for the kitchen magnets and larger for the industrial magnets. Also, when moving the two magnets apart, you can feel that the attractive forces and repulsive forces decrease. The mathematical relationship is described by Coulomb’s law: the force of attraction equals coulomb’s constant k x the quotient of Q1 x Q2 / r squared. The qs are the charges and r is the distance between the charges.
What does that mean for our atom? We have attractive forces between the nucleus and the electrons. The more protons in the nucleus, the stronger the attractive force. All electrons are experiencing the same attraction - there is no “splitting up”. It’s like love: you can love a lot of people with the same intensity. Sigh.
Back to Chemistry: Now, distance also matters: The greater the distance between the nucleus and the electron you are measuring, the lower the forces of attraction.
Let’s look at the periodic table: The number of protons increases across a period. Therefore, we can also say that the magnitude of attractive forces between nucleus and valence electrons is increasing.
If we go down a group, we are adding shells and are increasing the distance between the nucleus and the electrons and thus are decreasing the force between protons and valence electrons. Yes, down a group you are also increasing the number of protons, BUT, the increase in distance “matters more”.
Coulomb’s law provides us with a tool we can use to explain the periodic trends - properties that follow a pattern across the periods. These patterns help us to make predictions, for example about the reactivity of an element - and all without a crystal ball! And, Coulomb’s Law will help you in May to build your scientific argument when answering multiple choice questions!
We are focusing on four properties: atomic and ionic radii, ionization energy, electronegativity and electron affinity.
Atomic radius is defined as one-half the distance between the nuclei of identical atoms that are bonded together. Atomic radius decreases across a period, because the attractive force between nucleus and valence electrons increases. The electrons are pulled closer and therefore the atom is smaller. Atomic radius increases down a group, because you are adding shells aka energy levels and you are increasing the distance between the nucleus and the electrons.
To identify the patterns for ionic radii, we have to distinguish between cations and anions and compare them to their parent atom. Cations are generally smaller than their parent atom, because they have lost electrons on their highest energy level. As an example, the sodium cation is smaller than the sodium atom, because it has the electron configuration of neon.
Anions, on the other hand, are larger than their parent atoms. This is because they add electrons, which increases the electron-electron repulsion. I know, so far we have only talked about attractive forces, but the electrons are also repelling each other. Adding electrons increases the repulsion and the electron cloud spreads out. Therefore anions are larger than their parent atoms.
Ionization Energy, Electronegativity and electron affinity have similar trends and reasoning. So let’s start with the definitions and then combine them:
Ionization Energy is the energy needed to remove an electron - and form a cation. A similar, yet different, concept is electronegativity. It is the ability of an atom to attract a shared pair of electrons in a covalent bond. We’ve talked about this property earlier this season when recapping chemical bonding: It is the strength with which an atom can participate in a “tug of war”. Electron affinity is the “opposite” of ionization energy: electron affinity is the change in energy when an atom is accepting an electron and forming an anion. For many elements, the values for electron affinity are negative - which means energy is being released.
All three increase across a period due to the increase in nuclear charge, aka protons. With a higher number of protons we are increasing the attractive force between protons and valence electrons and are therefore increasing the energy needed to remove an electron, the pull on shared electrons and the energy released when adding an electron.
When we go down a group, all three of these properties decrease, because we are increasing the distance between the nucleus and the valence electrons and are therefore decreasing the attractive forces. Therefore, we have a lower ionization energy, lower electronegativity and release less energy when adding an electron.
There is a reason these properties are called trends: As it always is with a trend, there are exceptions. Here are some examples of anomalies you should be able to discuss: Oxygen has a lower first ionization energy than Nitrogen. Why? Oxygen has more protons! BUT, when you look at the electron configuration, you can see that nitrogen has 1 electron in each of the three p-orbitals. In oxygen, one of the p-orbitals has two electrons. This is like siblings sharing a room. Which leads to tension, uhm, increased electron-electron repulsion. And the sibling moves out more easily. Sorry!
Another exception is the first ionization energy of Boron and Beryllium. Using the trend, we would expect Boron to have a greater first ionization energy than Beryllium, but it actually takes more energy to remove the electron from Beryllium. The reason for this is that the valence electron in Beryllium is in the 2s orbital and for Boron the electron you would remove in the 2p orbital. The 2p orbital is already greater in energy - and therefore you need to add less energy yourself to remove the electron. The 2s orbital is lower in energy by itself, so you need to provide more energy to remove the electron.
We’ll end with an important test taking tip: The College Board will NOT accept a trend as an explanation. You will not receive points for stating that Fluorine is smaller than oxygen because atomic size decreases across a period. You HAVE TO explain the trend by using Coulomb’s Law!
To recap……
Coulomb’s Law states that the strength of the force is determined by the charge as well as the distance. The greater the charge and the shorter the distance, the higher the attractive forces.
Within a period the attractive forces generally increase due to an increase in nuclear charge. Down a group, the attractive forces generally decrease due to an increase in distance between nucleus and valence electrons. Atomic and ionic radii, ionization energy, electronegativity and electron affinity follow periodic trends that can be explained using Coulomb’s Law.
Coming up next on the APsolute RecAP Chemistry Edition: Chemical Bonds.
Today’s Question of the day is about Ionization energy.
Question: Removing a second electron and forming a 2+ cation, requires ________ energy than removing the first electron.